An ionic equation strips a full equation down to just the particles that actually change, cancelling spectator ions that are present on both sides unchanged. Mastering ionic equations shows an examiner you understand why a reaction occurs at the particle level, not just what the products are.
What is the difference between a full equation and an ionic equation?
A full (molecular) equation shows all reactants and products as complete formulae, including ions that do not participate in the reaction.
A full ionic equation splits all soluble ionic compounds and strong acids into their individual ions.
A net ionic equation (also called just the ionic equation at GCSE) cancels the spectator ions and shows only the particles that actually react.
Why it matters: two completely different reactions can have the same net ionic equation if the same type of chemical change takes place. Every neutralisation reaction between a strong acid and a strong alkali in solution produces the same ionic equation, regardless of which acid or alkali is used.
What are spectator ions?
Spectator ions are ions that appear on both the left and right sides of the full ionic equation, unchanged. They are present in solution throughout the reaction but do not form any new bonds or undergo any change. They are cancelled out when writing the net ionic equation.
A useful way to recognise them: if an ion appears with identical charge and formula on both sides, it is a spectator ion.
How to write an ionic equation: step-by-step method
Follow these four steps for any ionic equation:
- Write the balanced full equation using correct formulae.
- Identify which substances are ionic and soluble — these split into their ions. Solids, liquids, and gases are written as whole formulae (not split).
- Write the full ionic equation by splitting the soluble ionic compounds and strong acids/alkalis into individual ions. Keep the states in brackets: (aq), (s), (l), (g).
- Cancel spectator ions — cross out any ion that appears identically on both sides.
Worked example 1: neutralisation (strong acid + strong alkali)
Full equation: HCl(aq) + NaOH(aq) → NaCl(aq) + H₂O(l)
Full ionic equation (split all soluble ionic/acid species; keep water as H₂O): H⁺(aq) + Cl⁻(aq) + Na⁺(aq) + OH⁻(aq) → Na⁺(aq) + Cl⁻(aq) + H₂O(l)
Cancel spectator ions (Na⁺ and Cl⁻ appear on both sides): H⁺(aq) + OH⁻(aq) → H₂O(l)
This is the net ionic equation for every reaction between a strong acid and a strong alkali in solution. It is the same whether you use hydrochloric acid + sodium hydroxide, sulfuric acid + potassium hydroxide, or any other strong acid–strong alkali combination.
Worked example 2: precipitation reaction
Full equation: BaCl₂(aq) + Na₂SO₄(aq) → BaSO₄(s) + 2NaCl(aq)
The precipitate is BaSO₄, which is insoluble — write it as a whole formula.
Full ionic equation: Ba²⁺(aq) + 2Cl⁻(aq) + 2Na⁺(aq) + SO₄²⁻(aq) → BaSO₄(s) + 2Na⁺(aq) + 2Cl⁻(aq)
Cancel spectator ions (2Na⁺ and 2Cl⁻): Ba²⁺(aq) + SO₄²⁻(aq) → BaSO₄(s)
Any reaction that forms barium sulfate as a precipitate will have this same net ionic equation, regardless of which soluble barium salt and which soluble sulfate you start with.
Worked example 3: metal displacement
Full equation: Mg(s) + CuSO₄(aq) → MgSO₄(aq) + Cu(s)
Magnesium and copper are solids — written as whole formulae. CuSO₄ and MgSO₄ are soluble ionic compounds — split them.
Full ionic equation: Mg(s) + Cu²⁺(aq) + SO₄²⁻(aq) → Mg²⁺(aq) + SO₄²⁻(aq) + Cu(s)
Cancel spectator ions (SO₄²⁻): Mg(s) + Cu²⁺(aq) → Mg²⁺(aq) + Cu(s)
This shows clearly that magnesium atoms lose 2 electrons each (oxidation) and copper(II) ions gain 2 electrons each (reduction) — a redox reaction.
Worked example 4: acid with a carbonate
Full equation: CaCO₃(s) + 2HCl(aq) → CaCl₂(aq) + H₂O(l) + CO₂(g)
CaCO₃ is solid (insoluble) — write as whole formula. HCl splits into H⁺ and Cl⁻. CaCl₂ splits into Ca²⁺ and 2Cl⁻.
Full ionic equation: CaCO₃(s) + 2H⁺(aq) + 2Cl⁻(aq) → Ca²⁺(aq) + 2Cl⁻(aq) + H₂O(l) + CO₂(g)
Cancel spectator ions (2Cl⁻): CaCO₃(s) + 2H⁺(aq) → Ca²⁺(aq) + H₂O(l) + CO₂(g)
| Reaction type | Key indicator | Always true in net ionic equation |
|---|---|---|
| Strong acid + strong alkali | No precipitate, water formed | H⁺ + OH⁻ → H₂O |
| Precipitation | Insoluble product forms | The two ions forming the precipitate |
| Metal displacement | More reactive metal added to metal ion solution | Metal + ion → ion + metal |
| Acid + carbonate | Fizzing, CO₂ produced | Carbonate/solid + H⁺ → ion + H₂O + CO₂ |
Frequently asked questions
How do I know which compounds to split into ions?
Split ionic compounds that are soluble in water (aqueous, state symbol (aq)) and strong acids (HCl, H₂SO₄, HNO₃). Never split: insoluble solids (s), liquids (l), gases (g), weak acids such as ethanoic acid (CH₃COOH), or covalent molecules such as water. Using a solubility chart helps: common soluble ionic compounds include most sodium, potassium, and ammonium salts, most nitrates, and most chlorides except AgCl and PbCl₂.
Why do we bother with ionic equations if the full equation already tells us the products?
Ionic equations reveal the chemical mechanism more clearly — they show which species actually interact and which are just passengers in solution. They make it obvious that the same fundamental chemistry is happening across a whole family of reactions (all acid–alkali neutralisations in solution, for example). At a higher level, ionic equations also link directly to half-equations and electrochemistry, where tracking individual ions and electrons is essential.
What is a half-equation and how does it relate to ionic equations?
A half-equation shows what happens to just one species in a redox reaction — either the oxidation half or the reduction half, written separately. For the magnesium–copper(II) displacement above: oxidation half: Mg → Mg²⁺ + 2e⁻ and reduction half: Cu²⁺ + 2e⁻ → Cu. Adding them together (cancelling the 2e⁻) gives the net ionic equation. Half-equations are important in GCSE Higher tier for electrolysis questions.
What if I can't remember the solubility rules?
Exams usually provide or test a small set of rules. The safest rules to learn: all common sodium, potassium, and ammonium compounds are soluble; all nitrates are soluble; most common chlorides are soluble (except AgCl, PbCl₂); most sulfates are soluble (except BaSO₄, CaSO₄, PbSO₄); most carbonates and hydroxides are insoluble except those of sodium, potassium, and ammonium. When in doubt about state symbols, the question or your exam board's data sheet will usually provide them.
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